The enthalpy of combustion is one of the most frequently used energy terms in IB Chemistry. It appears in Topic 5 (Energetics) and plays a major role in calculating energy released by fuels, comparing fuel efficiency, and analyzing thermochemical equations. Understanding what the enthalpy of combustion means—and how to use it correctly—will help you tackle exam questions with confidence.
What Is the Enthalpy of Combustion?
The enthalpy of combustion (ΔHc°) is the enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions.
Key features:
- Complete combustion (not partial)
- One mole of the substance
- Standard conditions (298 K, 100 kPa)
- Substances in their standard states
This standardization ensures values are comparable across different substances.
The General Form of a Combustion Reaction
A typical combustion reaction is:
fuel + O₂ → CO₂ + H₂O
Examples:
1. Combustion of methane
CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l)
2. Combustion of ethanol
C₂H₅OH(l) + 3O₂(g) → 2CO₂(g) + 3H₂O(l)
3. Combustion of hydrogen
2H₂(g) + O₂(g) → 2H₂O(l)
Only complete combustion counts for ΔHc°.
Why Enthalpy of Combustion Is Always Negative
Combustion reactions are exothermic.
Energy is released because:
- Strong bonds form in the products (especially O–H and C=O bonds)
- The bonds formed are stronger than the bonds broken
As a result:
ΔHc° values are always negative, indicating energy is released to the surroundings.
For example:
- ΔHc°(CH₄) = –890 kJ/mol
- ΔHc°(ethanol) = –1367 kJ/mol
This is why fuels release heat when burned.
Measuring Enthalpy of Combustion (Calorimetry)
A calorimeter measures temperature change when a known mass of fuel burns.
Basic steps:
- Burn the fuel beneath a container of water.
- Measure the temperature change of the water.
- Use the equation:
q = mcΔT
Where:
- m = mass of water
- c = specific heat capacity
- ΔT = temperature change
- Use moles of fuel burned to find ΔHc.
However, simple school experiments often give lower-than-expected values because of heat loss.
Using Enthalpy of Combustion in Calculations
1. Comparing Fuels
Fuels with more negative ΔHc° values release more energy per mole.
Hydrocarbons generally show:
- Increasing ΔHc° with increasing chain length
- More carbon = more CO₂ and H₂O formed = more energy released
2. Hess’s Law Cycles
Enthalpies of combustion can be used to calculate enthalpies of formation or reaction.
For example:
ΔHreaction = ΣΔHc(reactants) – ΣΔHc(products)
This is commonly tested in longer IB calculation questions.
3. Determining Energy Density
Energy density compares fuels based on:
- Energy per gram
- Energy per volume
This relates to real-world applications such as fuel selection for transport.
Common IB Mistakes
- Forgetting that the combustion of hydrogen produces liquid water under standard conditions
- Using gaseous H₂O instead of liquid in equations
- Forgetting to calculate based on one mole of the substance
- Mixing up ΔHc° with ΔHf°
- Not balancing equations properly
- Not applying Hess’s law correctly
Being careful with physical states is especially important.
FAQs
Why does chain length increase enthalpy of combustion?
Longer hydrocarbons have more bonds, so burning them forms more stable products (CO₂ and H₂O), releasing more energy.
Why must combustion be complete?
Incomplete combustion forms CO or C (soot), which releases less energy. Standard values require complete oxidation.
Why is water liquid in combustion enthalpy definitions?
Under standard conditions (298 K), water’s standard state is liquid, not gas.
Conclusion
The enthalpy of combustion is the heat released when one mole of a substance is completely burned in oxygen under standard conditions. It is always negative because combustion is exothermic, and it is essential for comparing fuels, performing Hess’s law calculations, and understanding energetics in IB Chemistry. Mastering this concept strengthens your ability to analyze energy changes in chemical reactions.
